Kinetics and mechanism of quinolinium dichromate mediated oxidation of sugar alcohols in Bronsted acid media

Article abstract of DOI:10.1002/kin.21339

Bronsted acid catalyzed oxidation of certain sugar alcohols (polyols) has been studied by quinolinium dichromate (QDC) using aqueous sulfuric, perchloric, and hydrochloric acids at different temperatures. At constant acidity, reaction kinetics revealed the second-order kinetics with a first order in [Alcohol] and [QDC]. Zucker-Hammett, Bunnett, and Bunnett-Olsen criteria were used to analyze acid-dependent rate accelerations. Bunnett-Olsen plots of (log k + H_ν) versus (H_ν + log [H⁺]), and (log k) versus (H_ν + log [H⁺]) afforded slope values (? and ?^*, respectively)?>?0.47, suggesting that a water molecule acts as a prton transfer agent in the slow step of the mechanism in the oxidation of alcohols by QDC in the presence of aqueous sulfuric, perchloric, and hydrochloric acids.

Full text of DOI:10.1002/kin.21339

Received: 10 September 2019
DOI: 10.1002/kin.21339
Revised: 10 November 2019
Accepted: 26 November 2019
A R T I C L E
Kinetics and mechanism of quinolinium dichromate mediated
oxidation of sugar alcohols in Bronsted acid media
Satish Babu Kodali1,2
Narendar Reddy Jakku2
Chinna Rajanna Kamatala3,4
Rajeshwar Rao Yerraguntla2,3
1Department of Chemistry, Rayalaseema
University, Kurnool, India
Abstract
Bronsted acid catalyzed oxidation of certain sugar alcohols (polyols) has been studied
by quinolinium dichromate (QDC) using aqueous sulfuric, perchloric, and hydrochlo-
ric acids at different temperatures. At constant acidity, reaction kinetics revealed the
second-order kinetics with a first order in [Alcohol] and [QDC]. Zucker-Hammett,
Bunnett, and Bunnett-Olsen criteria were used to analyze acid-dependent rate accel-
2Department of Chemistry, Government City
College, Hyderabad, India
3Green Chemistry & Catalysis, Rajiv Gandhi
University of Knowledge Technologies,
Basar, India
4Department of Chemistry, Osmania
University, Hyderabad, India
+
erations. Bunnett-Olsen plots of (log k + Hꢀ) versus (Hꢀ + log [H ]), and (log k)
versus (Hꢀ + log [H ]) afforded slope values (ꢁ and ꢁ*, respectively) > 0.47, sug-
+
Correspondence
gesting that a water molecule acts as a prton transfer agent in the slow step of the
mechanism in the oxidation of alcohols by QDC in the presence of aqueous sulfuric,
perchloric, and hydrochloric acids.
Chinna Rajanna Kamatala, GreenChemistry&
Catalysis, RajivGandhiUniversityofKnowl-
edgeTechnologies, Basar, Nirmal, 504 107,
India.
Email: kcrajannaou@yahoo.com,
kcrchem@gmail.com
K E Y W O R D S
acidity functions, Bronsted acid catalysis, Bunnett-Olsen, Bunnett, oxidation of certain sugar alcohols,
quinolinium dichromate, Zucker-Hammett
compounds12–16 such as phenols, hydrocarbons, benzaldehy-
des, ethanolamines, and isoniazid, etc. In our earlier stud-
1
INTRODUCTION
During the past several decades, acid dichromate derived from
potassium dichromate in general,1–4 and Cr(VI) reagents with
heterocyclic moeity5–16 in particular, received the attention
of chemists all over the world due to their immense impor-
tance in synthetic chemistry. Dichromates of heterocyclic
bases such as pyridinium dichromate, imidazolium dichro-
mate, nicotinium dichromate, isonicotinium dichromate, and
quinolinium dichromate (QDC) were employed as oxidiz-
ing agents in organic synthesis.8–11 Almost all the dichro-
mate reagents were used for the selective oxidation of alco-
hols, because primary and secondary alcoholic groups present
in the amino acid residues of collagen is of immense bio-
logical importance. Balasubramanian and Pratibha were the
first to explore QDC as a synthetic reagent in the oxidation
of alcohols.11 However, over a period of time, QDC was
used for the selective oxidation of a wide range of organic
ies, we explored kinetic and mechanistic results of oxida-
tion of ethanolamines by QDC.16 On the other hand, bifunc-
tional molecules such as ꢂ-hydroxy ketones were used as
fine chemicals as well as building blocks for the synthesis of
several pharmaceutically important molecules, which found
their utility as antidepressants, HIV-protease inhibitors, and
antitumorals.17 These important features drove us to focus
our attention on a detailed kinetic study of oxidation of cer-
tain sugar alcohols (polyols, viz, ethylene glycol (EG), glyc-
erol, erythritol, xyletol, sorbitol, and mannitol) in the presence
of Bronsted acids (H2SO4, HCl, and HClO4) to explore the
mechanistic details (see Scheme 1). Our preliminary studies
revealed that reaction rates are faster in high acidic solutions.
In view of the observed significant acid catalysis, efforts were
made to analyze the observed kinetic results by linear free
energy relationships involving acidity functions.
Int J Chem Kinet. 2019;1–11.
wileyonlinelibrary.com/journal/kin
© 2019 Wiley Periodicals, Inc.
1

2
KODALI ET AL.
S C H E M E 1 Structure of different alcohols
was devided into two portions. A saturated solution of 2,4-
dinitophenylhydrazine was added to the first portion and was
kept aside for some time. The 2,4-dinltro phenyl hydrazone
derivative appeared as a precipitate, indicating that the end
product is aldehyde.18,19
2
EXPERIMENTAL
Ethanol, EG, glycerol, erythritol, xylitol, mannitol, and sor-
bitol were procured from Sigma Aldrich (India), Avra (India),
and SD-fine Chemicals (India) through a local vendor.
QDC was prepared according to the method of Balasum-
bramanian and Pratibha.11 To the aqueous solution of CrO3
(100 g/100 mL), 86 mL of quinoline was added in small
portions with constant stirring. The solution was diluted to
The second part of the obtained product was characterized
by the chromotropic acid test,20 because this test is one of the
most widely used methods for the identification and quantifi-
cation of different aldehydes in addition to the other analyt-
ical methods. A small amount of chromotropic acid reagent
was added to the product solution, followed by addition of
sulfuric acid in excess, and heated for a while. After cool-
ing the absorbance of resultant colored solution was mea-
sured by a UV-visible spectrophotometer, which indicated
absorption maximum in the range of 570-580 nm against the
reagent blank, which confirmed that the oxidation product is
aldehyde.
◦
400 mL with acetone and the contents cooled upto −20 C. An
orange color QDC solid separated out as a solid, which was
washed with acetone and recrystallised from water. Purity of
the dried compound was checked by elemental analysis and
infrared spectroscopic analysis. The IR spectra (recorded on
a spectrometer, Shimatzu, Japan) indicated the bands at 930,
875, 765, and 730 cm−1 characteristic of the dichromate ion.17
2.1 Kinetic studies
3
RESULTS AND DISCUSSION
Progress of the QDC-alcohol reaction was followed by
estimating the unreacted [QDC] by the spectrophotometric
method. The absorbance values were recorded on a double
beam Shimadzu UV-visible spectrophotometer equipped with
a thermostatic cell compartment. The kinetic data are repro-
ducible with a precision of 3% error.
3.1 Salient kinetic features
1. Under the conditions [Alcohol] ≫ [QDC], plots of ln
[A /At] versus time were found to be linear, passing
0
through origin indicating the first-order dependence in
[QDC], where A and At indicate absorbance values (opti-
cal density) at z0ero time and at any instant, respectively
(Figures 1–6).
2.2 Stoichiometry and products of oxidation
To determine the stoichiometry of the QDC-alcohol reac-
tion, QDC was taken excess over alcohol to ensure complete
conversion of alcohol to product. Absorbance of QDC was
recorded at every 1 h for about 1 day. From the change in
absorbance values of QDC, it was found that 1 mol of QDC
was needed to oxidize 1 mol of alcohol.
2. When [Alcohol] ≫ [QDC], the study revealed that plots
of k′ versus [Alcohol] were linear passing through origin.
Such plots depict the first-order kinetics in [Alcohol] (Fig-
ures 7–9 and Table 1).
3. A rate of the reaction was greatly influenced by a structural
change in the nucleus of alcohol.
Oxidation of ethanol gave acetaldehyde, whereas corre-
sponding carbonyl compounds (aldehydes) were produced in
the case of polyols. After completion of the reaction, solvent
ether was added to the reaction mixture. The resultant reac-
tion mixture was passed through a short-silicalgel column.
Ether was evaporated from the etherial-eluent, and the resul-
tant was concentrated futher. The concentrate thus obtained
4. Rate constants increased with temperature. Accordingly,
free energy of activation (∆G#) has been computed by
Eyring’s theory of reaction rates21–23 using a rate constant
(k) at any given temperature (T),
(ꢃ) = (ꢃt)(ꢄ∕Nh)exp(−Δꢅ#∕ꢄꢆ ).
(1)

KODALI ET AL.
3
1.5
1
0.35
0.3
0.25
0.2
0.15
0.1
0.05
0
y = 0.0196x
R² = 0.9992
y = 0.0033x
R² = 0.9978
0.5
0
0
10
20
30
40
50
60
Time (min)
0
20
40
60
80
100
F I G U R E 5 Plot of ln(A /At) versus time for the oxidation of
Time (min)
0
erythritol in the presence of HCl at [Erythritol] = 0.5 mol/dm3,
[HCl] = 0.167 mol/dm3, 103[QDC] = 1.33 mol/dm3, and 303 K
[Color figure can be viewed at wileyonlinelibrary.com]
F I G U R E 1 Plot of ln(A /At) versus time for the oxidation of
glycerol in the presence of H20SO4 at [Glycerol] = 0.083 mol/dm3,
[H2SO4] = 0.167 mol/dm3, 103[QDC] = 1.33 mol/dm3, and 303 K
[Color figure can be viewed at wileyonlinelibrary.com]
2
y = 0.0243x
R² = 0.9994
1
0
0
10
20
30
40
50
60
70
Time (min)
F I G U R E 6 Plot of ln(A /At) versus time for the oxidation of
sorbitol in the presence of H20SO4 at [Sorbitol] = 0.167 mol/dm3,
[H2SO4] = 0.167 mol/dm3, 103[QDC] = 1.33 mol/dm3, and 303 K
[Color figure can be viewed at wileyonlinelibrary.com]
F I G U R E 2 Plot of ln(A /At) versus time for the oxidation of
0
glycerol in the presence of HClO4 at [Glycerol] = 0.333 mol/dm3,
[HClO4] = 0.167 mol/dm3, 103[QDC] = 1.33 mol/dm3, and 303 K
[Color figure can be viewed at wileyonlinelibrary.com]
F I G U R E 7 Variation of pseudo–first-order rate constant (k′) with
[Ethanol] during the oxidation of ethanol in the presence of sulfuric
acid at [H2SO4] = 0.167 mol/dm3, 103[QDC] = 1.33 mol/dm3, and
303 K [Color figure can be viewed at wileyonlinelibrary.com]
F I G U R E 3 Plot of ln(A /At) versus time for the oxidation of
0
ethanol in the presence of H2SO4 at [Ethanol] = 0.167 mol/dm3,
[H2SO4] = 0.167 mol/dm3, 103[QDC] = 1.33 mol/dm3, and 303 K
[Color figure can be viewed at wileyonlinelibrary.com]
F I G U R E 8 Variation of pseudo–first-order rate constant (k′) with
[Erythritol] during the oxidation of erythritol in the presence of (A)
hydrochloric acid and (B) sulfuric acid medium at
[Acid] = 0.167 mol/dm3, 103[QDC] = 1.33 mol/dm3, and 303 K
[Color figure can be viewed at wileyonlinelibrary.com]
F I G U R E 4 Plot of ln(A /At) versus time at for the oxidation of
0
erythritol in the presence of HClO4 at [Erythritol] = 0.333 mol/dm3,
[HClO4] = 0.167 mol/dm3, 103[QDC] = 1.33 mol/dm3, and 303 K
[Color figure can be viewed at wileyonlinelibrary.com]
In the above equation, the terms R, N, h, and kt represent
gas constant, Avagadro number, Plank’s constant, and trans-
mission coefficient, respectively. But, the transmission coef-
ficient kt is generally equal to unity. Upon rearrangement of

4
KODALI ET AL.
F I G U R E 9 Variation of pseudo–first-order rate constant (k′) with
[Polyol] during the oxidation of (A) xylitol and (B) mannitol in the
presence of perchloric acid at [HClO4] = 0.167 mol/dm3,
103[QDC] = 1.33 mol/dm3, and 303 K [Color figure can be viewed at
wileyonlinelibrary.com]
F I G U R E 10 Variation of pseudo–first-order rate constant (k′)
with [Polyol] during the oxidation of sorbitol in the presence of sulfuric
acid at [H2SO4] = 0.167 mol/dm3, 103[QDC] = 1.33 mol/dm3, and
313K [Color figure can be viewed at wileyonlinelibrary.com]
Equation (1), followed by taking a natural logarithmic form
Equation (2) is obtained:
Rearrangment of Equation (4) afforded the following equa-
tion:
Δꢅ# = ꢄꢆ ln(ꢄꢆ ∕Nhk).
(2)
Δꢅ# = (−ꢆ Δꢈ# + Δꢇ#).
(5)
Substitution of R, N, and h values in SI units into Equa-
The enthalpy and entropies of activation (∆H# and ∆S#)
values were computed from the intercept and slopes of the
plots of ∆G# versus temperature (T). Few representative plots
are given in Figures 10–14. The activation parameters related
to various alcohols are presented in Table 2.
tion (2), ∆G# is simplified to Equation (3),
Δꢅ# = 8.314 × ꢆ [23.7641 + ln(ꢆ ∕ꢃ)].
(3)
Using Equation (3), ∆G# values were obtained at various
temeperatures. But, according to the Gibbs-Helmholtz equa-
tion (Equaton 4), ∆G# values could be used conveniently to
compute enthalpy and entropies of activation (∆H# and ∆S#)
values.
3.2 Acid catalysis and acidity functions
In the present study, the rate of the reaction was found to
depend on the nature and the concentration of the Bron-
sted acid (H2SO4, HCl, and HClO4). The rate of the
Δꢅ# = Δꢇ# − ꢆ Δꢈ#.
(4)
TA B L E 1 Concentration-dependent rate constants for the oxidation of alcohols by QDC in different acid media at 103[QDC] = 1.33 mol/dm3,
[Acid] = 0.167 mol/dm3, and 303K
Erythritol
k′ (min)
0.0014
0.0029
0.005
Xylitol
k′ (min)
0.002
0.006
0.011
0.026
0.040
0.004
0.007
0.014
0.034
0.048
0.006
0.012
0.032
0.063
0.111
Sorbitol
k′ (min)
0.0056
0.0074
0.019
0.041
0.065
0.004
0.009
0.019
0.050
0.072
0.006
0.009
0.024
0.042
0.079
Mannitol
k′ (min)
0.006
0.012
0.022
0.063
0.083
0.007
0.010
0.023
0.046
0.073
0.004
0.008
0.019
0.038
0.055
Acid
HClO4
[Alcohol] (mol/dm3)
0.042
(Rꢀ)
(Rꢀ)
(Rꢀ)
(Rꢀ)
0.965
0.973
0.994
0.999
0.999
0.989
0.977
0.998
0.996
0.999
0.944
0.989
0.988
0.991
0.998
0.991
0.979
0.999
0.998
0.995
0.931
0.971
0.999
0.999
0.999
0.996
0.994
0.993
0.995
0.995
0.947
0.996
0.991
0.986
0.987
0.994
0.993
0.997
0.982
0.982
0.999
0.982
0.999
0.991
0.996
0.998
0.998
0.997
0.998
0.999
0.991
0.983
0.999
0.998
0.990
0.993
0.996
0.998
0.999
0.996
0.083
0.167
0.333
0.011
0.500
0.017
HCl
0.042
0.002
0.083
0.003
0.167
0.006
0.333
0.013
0.500
0.020
H2SO4
0.042
0.004
0.083
0.008
0.167
0.015
0.333
0.033
0.500
0.047

KODALI ET AL.
5
+
calculation of [H ] in moderately and highly concen-
trated solutions could be expressed in terms of acidity
functions.23–29 Basically, Zucker and Hamett23,24 classified
acid-catalyzed reactions into A1 and A2 categories and pro-
posed following empirical relationships:
(1). In A1-type reactions, the rate-controlling step is uni-
molecular. The plots of log k versus Hammett’s acidity func-
tion −Hꢀ should be linear with unit slope (m = 1).
F I G U R E 11 Gibbs-Helmholtz plot for QDC oxidation of glycerol
in H2SO4 medium [Color figure can be viewed at
wileyonlinelibrary.com]
log ꢃ = 푚(ꢇꢀ) + log ꢃ푟.
(6)
Such linear plots with unit slope (m = 1) suggest that a
water molecule does not participate in the slow step.
84
83
(2). The A2-type reactions are solvolytic reactions in which
a water molecule is involved in the transformation of proto-
y = 0.131x + 40.577;
R2 = 0.9832
82
81
80
79
78
+
nated species (SH ) to the transition state. According to this
theory, the plots of log (rate constant) versus log [Acid] are
linear with unit slope (m# = 1), which is an indication of water
molecule participation in the bimolecular rate-determining
step.
290
295
300
305
310
T (K)
315
320
325
log ꢃ = 푚#log[H+] + log ꢃ푟.
(7)
F I G U R E 12 Gibbs-Helmholtz plot for QDC oxidation of
erythritol in HClO4 medium [Color figure can be viewed at
wileyonlinelibrary.com]
However, in the present study, the slopes of Zucker-
Hammett’s plots (m or m#) were found to be either greater
or less than unity, we thought that it is worthwhile to analyze
the rate data with Bunnett and Bunnett-Olsen’s theories.
3.2.1 Bunnett’s hypothesis
According to Bunnett’s hypothesis,24–27 the plots of [log (rate
constant + Hꢀ)] as a function of log 푎H Oshould be linear with
2
characteristic slopes of 휔.
(log ꢃ + Hꢀ) = 휔(log 푎water ) + 퐶.
(8)
F I G U R E 13 Gibbs-Helmholtz plot for QDC oxidation of mannitol
in HCl medium [Color figure can be viewed at wileyonlinelibrary.com]
On the other hand, the plot of [log (rate constant) – log
[acid]] as a function of log 푎ꢇ 푂should be linear with charac-
2
teristic slopes of 휔#.
82
81
y = 0.141x + 35.397
(log ꢃ − log[Acid]) = 휔∗(log 푎water ) + 퐶.
(9)
80
R² = 0.9894
79
78
77
76
On the basis of the magnitudes of observed slopes of (휔
and 휔#) the acid-catalyzed reactions were classified into three
categories.
290
300
310
320
330
T (K)
3.2.2 Bunnett-Olsen’s approach
F I G U R E 14 Gibbs-Helmholtz plot for QDC oxidation of sorbitol
in HClO4 medium [Color figure can be viewed at
wileyonlinelibrary.com]
To gain further insight into the observed trends of acid catal-
ysis, rate data have been cast into Bunnett-Olsen’s theories,28
which are applicable to the reactions involving weakly basic,
moderately basic, and strongly basic compounds as substrates.
(i) For weakly basic substrate, Bunnett-Olsen’s equation
was given as
oxidation reaction accelerated with increasing [acid] from 0.1
to 1.0 mol/dm3. In dilute acid solutions ([acid] < 0.1 mol/
+
dm3) [H ] can be easily expressed in terms of pH, but
it is not possible in more concentrated solutions ([acid]
≥ 0.10 mol/dm3). However, literature reports revealed that
(log ꢃ + ꢇꢀ) = 휑(ꢇꢀ + log[H+]) + log ꢃ푟.
(10)

6
KODALI ET AL.
TA B L E 2 Temperature-dependant rate constants and activation parameters for QDC oxidation of polyhydroxy compounds in H2SO4, HClO4,
and HCl media
Medium
Substrate
Ethanol
Temperature (◦K)
293
313
323
293
313
323
293
313
323
293
313
323
293
313
323
293
313
323
293
313
323
293
313
323
293
313
323
293
313
323
293
313
323
293
313
323
293
313
323
293
313
323
293
313
323
102 k″ (dm3/mol/s)
1.70
4.90
9.40
0.620
2.00
4.35
2.10
6.40
11.7
9.90
22.0
34.6
12.0
37.0
54.0
6.40
18.0
29.0
4.10
11.0
18.0
1.35
5.50
12.4
4.86
17.8
24.2
8.80
31.0
49.6
8.70
28.0
40.0
13.0
31.0
57.0
1.62
4.10
12.0
6.24
18.2
26.5
9.12
29.0
40.4
∆G# (kJ/mol)
81.7
84.6
85.7
84.1
86.9
87.7
81.1
83.9
85.1
77.3
80.7
82.2
76.9
79.4
81
∆H# (kJ/mol)
42.9
−∆S# (J/K/mol)
133
H2SO4
EG
47.5
41.2
29.3
37.0
36.5
36.45
55.7
40.6
43.3
37.7
35.4
42.0
35.4
36.7
125
136
164
136
143
147
91.0
131
117
136
141
136
147
Glycerol
Erythritol
Xylitol
Mannitol
Sorbitol
Glycerol
Erythritol
Xylitol
78.4
81.3
82.7
79.5
82.5
83.9
82.2
84.3
84.9
79.1
81.3
83.1
77.7
79.8
81.2
77.7
80.1
81.8
76.7
79.8
80.8
81.8
85.1
85.1
78.5
81.3
82.9
77.6
80.0
81.8
HClO4
Mannitol
Sorbitol
Glycerol
Erythritol
Xylitol
HCl
139
(Continues)

KODALI ET AL.
7
TA B L E 2 (Continued)
Medium
Substrate
Temperature (◦K)
102 k″ (dm3/mol/s)
∆G# (kJ/mol)
78.2
∆H# (kJ/mol)
41.5
−∆S# (J/K/mol)
125
Mannitol
293
313
323
293
313
323
6.95
23.0
38.0
9.80
28.0
51.0
80.6
81.9
Sorbitol
77.4
39.8
128
80.1
81.1
F I G U R E 15 Zucker-Hammet plot of ethylalcohol in QDC-H2SO4
F I G U R E 17 Bunnett plot of ethylalcohol in QDC-H2SO4 [Color
[Color figure can be viewed at wileyonlinelibrary.com]
figure can be viewed at wileyonlinelibrary.com]
F I G U R E 18 Bunnett plot of ethylalcohol in QDC-H2SO4 [Color
F I G U R E 16 Zucker-Hammet plot of ethylalcohol in QDC-H2SO4
figure can be viewed at wileyonlinelibrary.com]
[Color figure can be viewed at wileyonlinelibrary.com]
According to this equation plot of (log k + Hꢀ) versus
+
(Hꢀ + log [H ]) should be linear with a positive slope (휑)
and an intercept of log kr (specific rate constant).
(ii) For moderate and strongly basic substrates
log ꢃ = ꢁ#(ꢇꢀ + log[H+]) + log ꢃ푟.
(11)
According to this relationship plot of log k versus (Hꢀ +
+
log [H ]) should be linear with a positive slope (휑#) and a
F I G U R E 19 Bunnett-Olson plot of ethylalcohol in QDC-H2SO4
[Color figure can be viewed at wileyonlinelibrary.com]
definite intercept on ordinate (log kr).
In the present study, in each case, acid-dependent rate
constants were characterized by (a) Zucker-Hamett plots of
(log k versus Hꢀ and log k versus log [Acid]), (b) Bunnett plots
of [log (k + (Hꢀ)] versus log (aw), and [log k – log (acid)] ver-
sus log aw, and finally (c) Bunnett-Olsen plots of ((log k + Hꢀ)
+
+
versus (Hꢀ + log [H ]), and (log k) versus (Hꢀ + log [H ]).
Certain representative plots are given with their characteristic
slopes in Figures 15–20.
Data summarized in Table 3 show that the slopes of Zucker-
Hammett (m and m#), Bunnett (휔 and 휔#), and Bunnett-
Olsen (ꢁ or ꢁ*) plots have their own significance. These
F I G U R E 20 Bunnett-Olson plot of ethylalcohol in QDC-H2SO4
[Color figure can be viewed at wileyonlinelibrary.com]

8
KODALI ET AL.
TA B L E 3 Acidity function plots (QDC) in different acid media
H2SO4
Slope
HClO4
Slope
HCl
Slope
Compound
Theory/plot
Parameter
R2
R2
R2
Ethylene glycol
Zucker-Hammet I
Zucker-Hammet II
Bunnett III
m
1.056
0.927
0.988
0.988
0.947
0.981
0.998
0.994
0.981
0.874
0.897
0.991
0.998
0.970
0.986
0.933
0.921
0.960
0.990
0.998
0.946
0.876
0.859
0.989
0.999
0.973
0.987
0.922
0.902
0.969
0.989
0.965
0.996
0.938
0.930
0.979
0.989
0.945
0.995
0.946
0.942
0.971
0.984
1.559
0.979
0.968
0.909
0.954
0.971
0.989
0.984
0.971
0.914
0.940
0.996
0.991
0.999
0.945
0.879
0.908
0.998
0.999
0.997
0.960
0.908
0.925
0.998
0.998
0.990
0.911
0.850
0.920
0.974
0.997
0.979
0.968
0.909
0.954
0.971
0.989
0.984
0.971
0.914
0.940
0.996
0.991
0.989
0.809
0.875
0.927
0.890
0.779
0.948
0.996
0.924
0.913
0.930
0.999
0.998
0.999
0.893
0.866
0.914
0.994
0.999
0.806
0.814
0.905
0.999
0.694
0.919
0.995
0.862
0.819
0.899
0.979
0.999
0.809
0.875
0.927
0.890
0.779
0.948
0.996
0.924
0.913
0.930
0.999
0.998
m
−0.344
−0.062
−0.019
2.522
−0.197
−0.036
−0.022
0.742
−0.179
−0.038
−0.021
1.058
*
휔
Bunnett IV
휔
*
Bunnett-Olsen V
Bunnett-Olsen VI
Zucker-Hammet I
Zucker-Hammet II
Bunnett III
ɸ
ɸ
0.860
0.443
0.453
*
Ethyl alcohol
Glycerol
Erythritol
Xylitol
m
m
휔
1.098
0.546
0.738
−0.344
−0.058
−0.019
1.607
−0.184
−0.029
−0.017
0.600
−0.180
−0.028
−0.013
0.872
*
*
Bunnett IV
휔
Bunnett-Olsen V
Bunnett-Olsen VI
Zucker-Hammet I
Zucker-Hammet II
Bunnett III
ɸ
ɸ
0.505
0.338
0.411
*
m
m
휔
0.938
0.491
0.723
−0.298
−0.049
−0.018
1.282
−0.162
−0.027
−0.015
0.559
−0.175
−0.027
−0.013
0.851
*
*
Bunnett IV
휔
Bunnett-Olsen V
Bunnett-Olsen VI
Zucker-Hammet I
Zucker-Hammet II
Bunnett III
ɸ
ɸ
0.470
0.314
0.408
*
m
m
휔
0.687
0.553
0.886
−0.211
−0.027
−0.014
0.735
−0.184
−0.031
−0.017
0.647
−0.226
−0.026
−0.013
0.859
*
*
Bunnett IV
휔
Bunnett-Olsen V
Bunnett-Olsen VI
Zucker-Hammet I
Zucker-Hammet II
Bunnett III
ɸ
ɸ
0.393
0.340
0.455
*
m
m
휔
0.722
0.528
0.785
−0.205
−0.026
−0.012
0.882
−0.202
−0.037
−0.020
0.616
−0.185
−0.029
−0.014
0.922
*
*
Bunnett IV
휔
Bunnett-Olsen V
Bunnett-Olsen VI
Zucker-Hammet I
Zucker-Hammet II
Bunnett III
ɸ
ɸ
0.410
0.326
0.427
*
Mannitol
Sorbitol
m
m
휔
1.019
1.559
0.989
−0.326
−0.051
−0.018
1.354
−0.197
−0.036
−0.022
0.742
−0.179
−0.038
−0.021
1.058
*
*
Bunnett IV
휔
Bunnett-Olsen V
Bunnett-Olsen VI
Zucker-Hammet I
Zucker-Hammet II
Bunnett III
ɸ
ɸ
0.491
0.443
0.453
*
m
m
휔
1.030
0.546
0.738
−0.333
−0.054
−0.019
1.424
−0.184
−0.029
−0.017
0.600
−0.180
−0.028
−0.013
0.872
Bunnett IV
휔
*
Bunnett-Olsen V
Bunnett-Olsen VI
ɸ
ɸ
0.496
0.338
0.411
*

KODALI ET AL.
9
TA B L E 4 Zucker-Hammett, Bunnett, and Bunnett-Olsen’s
Since the rate of oxidation is found to be substantially cat-
alyzed by H2SO4, HClO4, and HCl, and a rate is dependant
on the type of anion present in these acids, it could be rea-
sonable to consider a protonated form of mineral acid (BH)-
hypothesis
Function of water
Slope of different
acidity functions
molecule in the rate
+
Slope value
determining step
bound QDC [(QH)CrO3BH] as an oxidizing species in the
Zucker-Hammett
(m) = 1.00
H2O molecule is not
involved
present study.
slope (m or m )
*
Further, the data presented in Table 3 suggest that Bunnett-
Olsen slopes are in the range of water molecule (H2O)—
acting as a proton trnsfer agent, which can be observed in
Table 4. Thus, on the basis of the foregoing discussion, we
have tried to propose a most plausible mechanism with the
participation of water molecule in the rate-limiting step as a
proton transfer reagent as shown in Scheme 2. For the above
mechanism given in Scheme 2, the rate law could be derived
as
(m) ≠ 1.00
−2.5 to 0.0
+1.2 to 3.3
+3.3 to 7.0
< (−2)
H2O molecule may be
involved
Bunnett’s slope
H2O not involved
(휔)
H2O acts as a
nucleophilic agent
H2O acts as a proton
abstracting agent
Bunnett’s slope
H2O acts as a
−푑[QDC]
(휔 )
nucleophilic agent
= ꢃ′′[QDC][ROH]ꢇ푛
*
푑푡
> (−2)
H2O acts as a proton
abstracting agent
At constant acidity and [salt], the rate law could be simpli-
Bunnett-Olsen’s
−0.34 to 0
+0.18 to 0.47
> 0.47
H2O molecule is not
involved
fied as
slope (휑 or 휑#)
−푑[QDC]
H2O acts as a
= ꢃ′′[QDC][ROH]
nucleophilic agent
푑푡
H2O acts as a proton
transfer agent
The rate law is in consonance with the observed reac-
tion kinetics, that is, the first-order dependence on both the
reactants [QDC] as well as [Alcohol] at constant acidity and
ionic strength (휇). Presented rate constants, enthalpy activa-
tion (∆H#), and entropy of activation (∆S#) in Table 2 could be
summarized to follow the as shown below, eventhough there
is no consistency from acid to the other:
values could be interpreted by important criterion (features) of
Zucker-Hammett, Bunnett, and Bunnett-Olsen’s theories are
compiled in Table 4, which could be useful to ascertain the
involvement of water molecule in the rate-limition step and
thereby the its role to understand the type of mechanism. In
almost all the studied alcohols, Bunnett’s (휔#) and Bunnett-
Olsen’s (휑*) values are in favour of water molecule participa-
tion in the rate-limiting step as a proton abstracting reagent.
General trends of rate constants (k) in Table 2
H2SO4
HClO4
HCl
xylitol > mannitol > sorbitol > erythritol >
glycerol > EtOH > EG
sorbitol > = mannitol > = xylitol > erythritol >
glycerol
3.3 Reactive species and mechanism
sorbitol > = xylitol > mannitol > erythritol > glycerol
Presented trends in enthalpy of activation (∆H#) in Table 2
It was earlier been reported that Cr(VI) exists in several reac-
−
+
tive forms such as HCrO4 , H2CrO4, [HCrO3] , and HCrO3B
H2SO4
erythritol < xylitol < mannitol = sorbitol <
EtOH < EG
−
−
(where B = HSO , ClO , Cl ) in aqueous acidic solutions.1,2
Since QDC rese4mbles 4K2Cr2O7, a similar type of Cr(VI)
species could be proposed to exist in aqueous acidic solutions
according to the following euilibria:
−
HClO4
HCl
sorbitol < mannitol < xylitol < erythritol < glycerol
erythritol < xylitol < sorbitol < mannitol < glycerol
Presented trends in entropy of activation (∆S#) in Table 2:
H2SO4
erythritol < sorbitol < mannitol < xylitol =
glycerol < EtOH < EG
(QH)2Cr2O7 + H2O ⇌ 2H(QH)CrO4
HClO4
HCl
sorbitol < mannitol < erythritol < xylitol < glycerol
mannitol < sorbitol < glycerol < xylitol < erythritol
(where Q is quinoline and QH is quinolinium ion)
H(QH)CrO4 + H+ ⇌ (QH)CrO+ + H2O
3
Foregoing trends could not be correlated quantitavely
because of irregular trends. This may probably indicate the
cumulative effects of several parameters such as polar, steric,
hydrophobic, and hydrophic interactions, which control the
reaction rates. However, observed large negative entropies of
(QH)CrO + BH ⇌ [(QH)CrO3BH]+
(where B = HSO−4 , OAc−, Cl−, ClO and BH = H2SO4,
HOAc, HCl, HClO4)

10
KODALI ET AL.
S C H E M E 2 Bronstead acid dependant QDC
O
Cr
B
oxidation of alcohols
O-QH
OH
H
O
O-QH
OH
OH
+
O
+
Cr
H
R
B
R
(I) Bronsted acid bound QDC
(BHQDC)
Alcohol
H2O
H3O+
R
O
Slow (k)
Aldehyde
O-QH
Fast
O
H
O
O-QH
OH
B
Cr
Cr
OH
OH
[ (QH)OCrB(OH)2 ] or Cr(IV)
B
R
Here R- represents
Ethanol = -CH3 ; Glycol = - CH2OH
Glycerol = - CHOH(CH2OH) ;Erythritol = - (CHOH) 2 CH2OH
Xylitol = - (CHOH)3CH2OH ; Mannitol = -(CHOH) 4 CH2OH
Sorbitol = - (CHOH) 4CH2OH
Here [(QH)OCrB(OH)2 ] = Cr(IV)]; Q = Quinoline; QH= Quinolinium; B = ClO4 , HSO4 and Cl
activation (∆S#) probably suggest a systematic arrangement of
atoms in the transistion state due to solvation; and the resulted
activated complex may be polar than the reactants in terms of
the transition state. Further, the overall entropy change (∆S#)
accompanying the formation of the activated complex may
suggest that there is a loss in the transitional energy.13–15 The
∆G# values for all the reactions are by and large appeared to
be constant, which may probably indicate that a similar type
of mechanism is operative in the present reaction series.
quantitavely with any specific parameter because of irregular
trends. However, these obsrvations may probably suggest the
cumulative effects of several parameters such as polar, steric,
hydrophobic, and hydrophic interactions, which control the
reaction rates. Observed large negative entropies of activation
(∆S#) probably suggest a systematic arrangement of atoms in
the transistion state due to solvation; and the resulted activated
complex may be polar than the reactants in terms of the transi-
tion state. Further, the overall entropy change (∆S#) accompa-
nying the formation of the activated complex may suggest that
there is a loss in the transitional energy.21–23 The ∆G# values
for all the reactions are by and large appeared to be constant,
which may probably indicate that a similar type of mechanism
is operative in the present reaction series.
4
CONCLUSIONS
The QDC-alcohol reaction obeyed the second-order kinetics
with a first-order dependence on [alcohol] and [QDC] at con-
stant acidity. The rate of oxidation was found to be catalyzed
by mineral acids, namely HClO4, H2SO4, and HCl, etc. Rate
data analysis using Bunnett-Olsen’s theory of acidity func-
tions revealed that the plots of (log k + Hꢀ) versus (Hꢀ + log
ACKNOWLEDGMENTS
We greatly acknowledge the authorities of Osmania Univer-
sity Hyderabad for laboratory facilities, and Rayalaseema
University, Kurnool, India, for constant encouragement
extended to Mr. Kodali Satishbabu (Registration No: PP-
028), who has registered for Ph.D. proagram. Authors are
highly grateful to Professor P.K. Saiprakash (former Dean,
Science Faculty, O.U.) for constant encouragement.
+
+
[H ]), and (log k) versus (Hꢀ + log [H ]) depicted slope val-
ues (ꢁ and ꢁ*, respectively) > 0.47. These observations sug-
gested that a water molecule acts as a proton transfer agent in
the slow step of the mechanism in the oxidation of alcohols by
QDC in all the acid media, which were used in this study. But
the observed trends of rate constants could not be correlated

KODALI ET AL.
11
CONFLICT OF INTEREST
17. Miller FA, Wilkins CH. Infrared spectra and characteristic frequen-
cies of inorganic ions. Anal Chem. 1952;24:1253-1294.
18. Vogel AI. Elementary Practical Organic Chemistry (Part III):
Quantitative Oryanic Analysis. London: ELBS; 1971:735.
19. Dey BB, Raman MVS. Laboratory Manual of Organic Chem-
istry. Madras, India: Central Art Press; 1957:359 (revised by T. R.
Govinda Chary).
Authors declare that they do not have any conflict of interest.
ORCID
Chinna Rajanna Kamatala
https://orcid.org/0000-0003-0473-563X
20. (a) Feigl F, Anger V. Spot Tests in Inorganic Analysis. Amsterdam,
The Netherlands: Elsevier Science; 1998. (b) Parise Georghioua PE,
Ho CK. The chemistry of the chromotropic acid method for the anal-
ysis of formaldehyde. Can J Chem. 1989;67:871-876. (c) Dar A,
Shafique U, Anwar J, Waheed-uz-Zaman NaseerA. A simple spot
test quantification method to determine formaldehyde in aqueous
samples. J Saudi Chem Soc. 2016;20:S352-S356. (d) Fagnani E,
Melios CB, Pezza L, Pezza HR. Chromotropic acid-formaldehyde
reaction in strongly acidic media. The role of dissolved oxygen and
replacement of concentrated sulphuric acid. Talanta. 2003;60:171-
176.
REFERENCES
1. Wiberg KB. Oxidations in Organic Chemistry (Part-A). New York,
NY: Academic Press; 1965.
2. Benson D. Mechanisms of Oxidatioin by Metal ion. New York, NY:
Elsevier Scientific Publishing; 1976.
3. Fieser M, Fieser LF. Reacgents for Organic Synthesis. New York,
NY: Wiley-Interscience; 1977. Vol. 6.
4. Collins JC, Hess WW, Franck FJ. Dipyridine-chromium(VI)
oxide oxidation of alcohols in dichloromethane. Tetrahedron Lett.
1968;30:3363-3366.
21. Glasstone S, Laidler KJ, Eyring H. Theory of Rate Processes. New
York, NY: McGraw-Hill; 1961.
22. Espenson JH. Chemical Kinetics and Reaction Mechanisms. New
York, NY: McGraw-Hill; 2002:156-160.
5. Patel S, Mishra BK. Chromium(VI) oxidants having quaternary
ammonium ions: studies on synthetic applications and oxidation
kinetics. Tetrahedron. 2007;63:4367-4406.
23. Laidler KJ. Chemical Kinetics. Singapore: Pearson Education;
2004.
6. Karunakaran C, Karuthapandian S, Suresh S. Kinetic evidence of a
common mechanism in the oxidations of diethyl sulfide by dichro-
mates and halochromates of heterocyclic bases. Int J Chem Kinet.
2003;35:1-8.
24. (a) Zucker L, Hammett LP. The mechanism of the acid cat-
alyzed enolization of acetophenone derivatives. J Am Chem Soc.
1939;61:2785-2279. (b) Zucker L, Hammett LP. Kinetics of the iod-
ination of acetophenone in sulfuric and perchloric acid solutions. J
Am Chem Soc. 1939;61:2791-2798.
7. Karunakaran C, Suresh S. Similar substituent effects in the oxida-
tions of primary aliphatic alcohols with dichromates and halochro-
mates of heterocyclic bases. Int J Chem Kinet. 2005;37:5-9.
8. Corey EJ, Schmidt G. Useful procedures for the oxidation of alco-
hols involving pyridinium dichromate in aprotic media. Tetrahedron
Lett. 1979;20:399-402.
25. Paul MA, Long FA. H0 and related indicator acidity function. Chem
Rev. 1957;57:1-45.
26. Hammett LP. Physical Organic Chemistry. Tokyo, Japan: McGraw-
Hill, Kogakusa; 1970.
27. Rochester HS. Acidity Functions. New York, NY: Academic Press;
1970.
9. Kim S, Lhim DC. Imidazolium dichromate. A new reagent for the
oxidation of alcohols to carbonyl compounds. Bull Chem Soc Jpn.
1986;59:3297-3298.
28. (a) Bunnett JP. Kinetics of reactions in moderately concentrated
aqueous acids. I. classification of reactions. J Am Chem Soc.
1961;83:4956-4967. (b) Bunnett JP. Kinetics of reactions in mod-
erately concentrated aqueous acids. II. An empirical criterion of
mechanism. J Am Chem Soc. 1961;83:4968-4973. (c) Bunnett JP.
Kinetics of reactions in moderately concentrated aqueous acids. III.
Theory of w- and w*-values. J Am Chem Soc. 1961;83:4973-4977.
29. (a) Bunnett JP, Olsen FP. Linear free energy relationships con-
cerning equilibria in moderately concentrated mineral acids a sim-
ple method for estimating pK’s of weak bases. Can J Chem.
1966;44:1899-1916. (b) Bunnett JP, Olsen FP. Linear free energy
relationships concerning reaction rates in moderately concentrated
mineral acids. Can J Chem. 1966;44:1917-1931.
10. Lopez C, Gonzalez A, Cossio FP, Palomo C. 3-Carboxypyridinium
dichromate (NDC) and (4-carboxypyridinium dichromate (INDC).
Two new mild, stable, efficient, and inexpensive chromium(VI) oxi-
dation reagents. Synth Commun. 1985;15:1197-1211.
11. Balasubramanian R, Prathiba V.Quinolinium dichromate; a new
reagent for oxidation of alcohols. Indian J Chem B. 1986;25:326.
12. Aruna K, Manikyamba P. Linear free energy relationships in the
kinetic study of oxidation of phenols by quinolinium dichromate.
Int J Chem Kinet. 1997;29:437-443.
13. Raviraj MK, Dinesh CB, Nandibewoor ST. Oxidation of isoniazid
by quinolinium dichromate in an aqueous acid medium and kinetic
determination of isoniazid in pure and pharmaceutical formulations.
Anal Sci. 2004;20:743-747.
14. Medien HAA. Kinetics of oxidation of benzaldehydes by quinolin-
ium dichromate. Z Naturforsch b. 2003;58:1201-1205.
15. Sarma GC, Mahanti MK. Kinetics of oxidation of hydrocarbons by
quinolinium dichromate. J Phys Org Chem. 1991;4:217-224.
16. Narendar Reddy J, Giridhar S, Rajanna KC. Equilibrium and kinetic
studies of electron transfer reactions involving quinolinium dichro-
mate and aminoalcohols in aqueous acid media. Transition Met
Chem. 1996;21:105-111.
How to cite this article: Kodali SB, Jakku NR,
Kamatala CR, Yerraguntla RR. Kinetics and mecha-
nism of quinolinium dichromate mediated oxidation
of sugar alcohols in Bronsted acid media. Int J Chem
Kinet. 2019;1-11. https://doi.org/10.1002/kin.21339